ASV_ Unit 4 Chemical Reactions
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Unit 4 Chemical Reactions Student led notes with practice template
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Topic 4.1 Introduction to Reactions (MH3.7)
Chemical equations represent substances that change properties or transform into different substances
Objective: Identify evidence of chemical and physical changes in matter
Physical changes: change in properties but not composition
Chemical changes: transformation into new substances with different compositions
Possible evidence of chemical change: production of heat or light, formation of a gas or precipitate, color change
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4.1 Introduction to Reactions- Aryn
Chemical reaction: substance(s) change into another substance
Chemical equation: chemical symbols represent what happens during a chemical reaction
Reactants: starting materials
Products: ending materials
Symbols (g), (l), (s), (aq) indicate states of matter
Other symbols used in chemical equations: +- (separates reactants/products), ------> (yields), heat (required for reaction), O (catalyst present)
Practice Problem: Li + H3PO4 ---> H2 + Li3PO4
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Topic 4.2 Net Ionic Equations (MH 4.2)
Learning Objectives: Represent changes in matter with a balanced chemical or net ionic equation
Different forms of equations: balanced molecular, complete ionic, and net ionic equations
Form used depends on the context
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Molecular Equations
Equations where compounds are written as molecules or whole units
Does not show detailed changes
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Ionic Equations
Equations where species dissolved in water are shown as free ions
Spectator ions are not involved in the overall reaction
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Net Ionic Equations
Equations that only show the species involved in the reaction
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Topic 4.3 Representation of Reactions (MH3.7, 3.8)
Chemical equations represent chemical changes
Equations must contain equal numbers of atoms of every element before and after the change
Mass is conserved in chemical reactions
Balanced chemical equations can be translated into symbolic particulate representations
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Chemical Equations
Chemical reactions are represented by chemical equations
Symbols represent reactants, products, and states of matter
Equations must be balanced
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Law of Conservation of Mass
Matter cannot be created or destroyed
Chemical equations should have equal amounts of each element before and after the reaction
Balance equations according to subscripts and coefficients
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Balancing Chemical Equations
Identify reactants and products and write their formulas on the correct side of the equation
Try different coefficients to make the number of atoms of each element the same on both sides
Balance elements that appear only once on each side
Balance elements that appear in multiple chemical formulas on the same side
Check equation to ensure the same total number of each type of atom on both sides
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Particulate Representations
Chemical reactions can be presented as particulate representations
Depict atoms before and after the reaction to show changes
Representation should have equal numbers of atoms before and after the reaction
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Practice Problems
Matter cannot be created or destroyed
Balance given equations
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Draw a particulate representation for the Balanced Chemical Equation
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Topic 4.4 Physical and Chemical Changes (MH 11.8)
Objective: Explain the relationship between macroscopic characteristics and bond interactions for chemical and physical processes
Chemical processes involve breaking and/or formation of chemical bonds
Physical processes involve changes in intermolecular interactions, such as phase changes
Physical processes can involve breaking of chemical bonds
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Physical changes
Changes in intermolecular forces
Do not affect chemical composition or bond breaking/creating
Examples: water evaporating, slicing a loaf of bread
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Chemical changes
Involve breaking/creating of bonds
Change chemical composition
Examples: log combusting, boiling eggs, digestion
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Exception
Bonds can be broken during a physical change
Dissociation of ions in water as an example
Reactions are reversible
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Example problems
Identify physical and chemical changes
Explain why a candle involves both physical and chemical changes
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Heating Curve
Observations during heating of a solid below its melting point
Temperature increases at a constant rate until melting begins
Temperature remains constant during melting
Temperature increases at a constant rate until boiling begins
Temperature remains constant during boiling
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Heating Curve cont..
Energy added during temperature increase can be calculated using q=m c ΔT
q: energy, m: mass, c: specific heat capacity, ΔT: temperature change
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Phase Diagrams
Graph describing conditions of temperature and pressure for solid, liquid, and gas states
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Phase Diagrams
Melting point of water changes with increasing pressure
Stoichiometry (Page 24)
Atoms must be conserved during a chemical process
Balanced chemical equations contain information about the proportionality of substances
Coefficients can be used in chemical calculations involving the mole concept
Stoichiometric calculations can be combined with the ideal gas law and calculations involving molarity to study gases and solutions quantitatively
Stoichiometry (Page 25)
Stoichiometry is the quantitative study of reactants and products in a chemical reaction
It answers the question of how much product will be formed from a reaction with a specific amount of reactants
The mole method is used to calculate the amounts of substances involved in the reaction
Stoichiometry pt 2 (Page 26)
Conversion factors are used in stoichiometric calculations
The concept of stoichiometric equivalence is important for dimensional analysis
Different mole ratios can be used in calculations
Stoichiometry pt 3 (Page 27)
Steps in stoichiometry calculations:
Write a balanced equation for the reaction
Convert the given amount of the reactant to moles
Use the mole ratio from the balanced equation to calculate the moles of product formed
Convert the moles of product to grams or other units
Stoichiometry SS (Page 28)
Mass and molar mass are used in stoichiometry calculations
Mole ratios are used to relate the amounts of substances in a balanced equation
Stoichiometry practice problems (Page 29)
Examples of stoichiometry problems:
Calculating the mass of CO2 produced from a given amount of C6H12O6
Calculating the grams of Li needed to produce a given amount of H2
Introduction to Titration (Page 30)
Titration is used to determine the concentration of an analyte in solution
The equivalence point is when the analyte is completely consumed by the titrant
The endpoint is an observable event that indicates the equivalence point
Acid-Base Titration (Page 31)
Acid-base titration involves adding a standard solution of known concentration to an unknown solution until the reaction is complete
The concentration of the unknown solution can be calculated using the volumes and concentrations of the solutions used
A pH indicator is used to monitor the progress of the reaction
Example of Titration (Page 32)
Sodium hydroxide is standardized by titrating it against a known concentration of an acid solution
The equivalence point is reached when the acid has completely reacted with the base
Phenolphthalein is a commonly used indicator in acid-base titrations
Practice Problem of Titration (Page 33)
A student neutralizes a known amount of KHP with a NaOH solution
The concentration of the NaOH solution can be calculated using the volumes and masses of the solutions used
Redox Titration (Page 34)
Redox reactions involve the transfer of electrons
Oxidizing agents can be titrated against reducing agents
Indicators are used to monitor the progress of the reaction
Redox Titration cont. (Page 35)
Common oxidizing agents are KMnO4 and K2Cr2O7
Calculations in redox titrations are similar to acid-base titrations but slightly more complex due to stoichiometry
Practice Problem of Redox Titration (Page 36)
A volume of KMnO4 solution is needed to oxidize a FeSO4 solution in an acidic medium
The concentration of the FeSO4 solution can be calculated using the volumes and concentrations of the solutions used
Types of Chemical Reactions (Page 37)
Acid-base reactions involve the transfer of protons
Oxidation-reduction reactions involve the transfer of electrons
Precipitation reactions involve the formation of insoluble or sparingly soluble compounds
Types of Chemical Reactions continued... (Page 38)
Oxidation numbers can be assigned to identify the oxidized and reduced species in a redox reaction
Precipitation reactions involve mixing ions in aqueous solution to produce insoluble compounds
Solubility rules can be used to determine the solubility of salts in water
Types of Chemical Reactions (Page 39)
Reactions can be classified as acid-base, oxidation-reduction, or precipitation reactions
Page 40: Precipitation Reactions
Precipitation reactions result in the formation of an insoluble product, or precipitate.
A precipitate is an insoluble solid that separates from the solution.
Precipitation reactions usually involve ionic compounds.
Example: Pb(NO3)2 (aq) + 2KI (aq) —> PbI2 (s) + 2KNO3 (aq)
This type of reaction is also called a metathesis reaction or a double-displacement reaction.
Page 41: Solubility
Solubility refers to the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.
Solubility can predict whether a precipitate will form when a compound is added to a solution or when two solutions are mixed.
There are solubility rules that can be used to determine the solubility of common ionic compounds in water at 25°C.
Page 42: Solubility Rules
Table 4.2 provides solubility rules for common ionic compounds in water at 25°C.
Soluble compounds include those containing alkali metal ions (Li+, Na+, K+, Rb+, Cs+) and the ammonium ion (NH4+).
Exceptions include compounds containing certain ions such as nitrates (NO3-), acetates (CH3COO-), bicarbonates (HCO3-), chlorates (ClO3-), and perchlorates (ClO4-).
Insoluble compounds include carbonates (CO3^2-), phosphates (PO4^3-), chromates (CrO4^2-), sulfides (S^2-), and hydroxides (OH-).
Exceptions include compounds containing certain ions such as sulfates (SO4^2-) of Ag+, Ca2+, Sr2+, Ba2+, Hg2+, and Pb2+.
Page 43: Example Problem
The following ionic compounds are classified as soluble or insoluble:
(a) silver sulfate (Ag2SO4): Insoluble
(b) calcium carbonate (CaCO3): Insoluble
(c) sodium phosphate (Na3PO4): Soluble
The classification is based on the solubility rules and the type of ions present in the compounds.
Page 44: Molecular Equations, Ionic Equations, and Net Ionic Equations
The equation Pb(NO3)2 (aq) + 2KI (aq) —> PbI2 (s) + 2KNO3 (aq) is a molecular equation.
In a molecular equation, the formulas of the compounds are written as though all species existed as molecules or whole units.
An ionic equation shows dissolved species as free ions.
The net ionic equation shows only the species that actually take part in the reaction.
Page 45: Procedure for Writing Ionic and Net Ionic Equations
The procedure for writing ionic and net ionic equations involves the following steps:
Write a balanced molecular equation for the reaction using the correct formulas for the reactant and product ionic compounds.
Refer to Table 4.2 to determine which of the products is insoluble and will appear as a precipitate.
Write the ionic equation for the reaction, showing the compound that does not appear as the precipitate as free ions.
Identify and cancel the spectator ions on both sides of the equation to write the net ionic equation.
Check that the charges and number of atoms balance in the net ionic equation.
Page 46: Acid Base Part 1
Acids produce H+ ions when added to water through ionization.
They have a sour taste and can change the color of plant dyes.
Acids can react with specific metals like zinc, magnesium, and iron to produce hydrogen ions.
They can also react with carbonates or bicarbonates to produce carbon dioxide gas.
Acids conduct electricity in their aqueous state.
Bases produce OH- ions when added to water.
They have a bitter taste and can change the color of plant dyes.
Bases have a smooth/slippery feeling, similar to soap or oil.
Bases also conduct electricity in their aqueous state.
Page 47: Acid Base Part 2
Brønsted acids donate protons, while Brønsted bases accept protons.
The hydronium ion (H3O+) is formed when water gains an extra H ion.
Monoprotic acids yield one hydrogen ion upon ionization.
Diprotic acids yield two hydrogen ions, and triprotic acids yield three hydrogen ions.
Page 48: Acid Base Part 3
Double replacement reactions occur when an acid-base reaction takes place, and the positive/negative ions switch places to produce two new substances.
Page 49: Acid Base Example
Identify whether the bolded compounds are Bronsted bases or Bronsted acids:
HSO4- + NH3 → NH4+ + SO4-: HSO4- is a Bronsted acid, NH3 is a Bronsted base.
HCl + HSO4- → H2SO4 + Cl-: HCl is a Bronsted acid, HSO4- is a Bronsted base.
Identify the acids as monoprotic, diprotic, or triprotic:
HNO3(aq): Monoprotic
H2SO4: Diprotic
H2CO3: Monoprotic
H3PO4(aq): Triprotic
Page 50: Oxidation Reduction
Oxidation-reduction (redox) reactions involve the transfer of one or more electrons between chemical species.
Combustion is a subclass of redox reactions where a species reacts with oxygen gas.
Redox reactions are also known as electron-transfer reactions and are used in obtaining metallic and nonmetallic elements from their ores.
Page 51: Example of Oxidation Reduction (Redox Reaction)
The reaction between oxygen and magnesium can be represented by the formula 2Mg + O2 → 2MgO.
Each step of the reaction can be shown as a half-reaction, explicitly indicating the gain and loss of electrons.
The two half-reactions can be combined to create the full redox reaction, with the electrons canceled on both sides of the equation.
Page 52: Redox Reaction Terminology
The oxidation reaction is the half-reaction that involves the loss of electrons.
The reduction reaction is the half-reaction that involves the gain of electrons.
The reducing agent is the element that donates electrons to other elements and causes them to be reduced.
The oxidation agent is the element that accepts electrons from other elements.
Page 53: Practice Problem
Show each step of the redox reaction for the following reaction: Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)
The steps involve writing the half-reactions for the oxidation and reduction, and then combining them to form the full redox reaction.
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Oxidation Number
Charge an atom would have if its electrons were transferred completely
Despite molecular compounds not actually transferring electrons, oxidation numbers can still be used and redox reactions can still be calculated
Experimental measurements show there are partial transfers of electrons
Example: H₂ (g) + Cl₂ (g) → 2HCl (g)
H₂ and Cl₂ have 0 as the oxidation number
In the HCl molecule, H has a 1+ and Cl has a 1-
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Oxidation number rules
Free elements have an oxidation number of zero
Example: H₂, Br₂, Na, Be, K, O₂, P₄
For monatomic ions, the oxidation number is equal to the charge on the ion
Example: Li+ ion has an oxidation number of +1, Ba₂+ ion has +2, I− ion has −1, O₂− ion has −2
All alkali metals have an oxidation number of +1 and all alkaline earth metals have an oxidation number of +2 in their compounds
Aluminum has an oxidation number of +3 in all its compounds
The oxidation number of oxygen in most compounds is −2, but in hydrogen peroxide (H₂O₂) and peroxide ion (O₂²⁻), it is −1
The oxidation number of hydrogen is +1, except when it is bonded to metals in binary compounds, then its oxidation number is −1
Fluorine has an oxidation number of −1 in all its compounds
Other halogens (Cl, Br, and I) have negative oxidation numbers unless combined with oxygen, then they have positive oxidation numbers
In a neutral molecule, the sum of the oxidation numbers must be zero
In a polyatomic ion, the sum of oxidation numbers of all the elements in the ion must be equal to the net charge of the ion
Example: NH₄+, N is −3 and H is +1
Oxidation numbers do not have to be integers
Example: The oxidation number of O in the superoxide ion, O₂−, is −1/2
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Practice Problems
Assign oxidation numbers to all the following compounds/ions
PF₃: P= +3, F= -1
MnO₄⁻: Mn= +7, O= -2
KO₂: K= +1, O= -1/2
IF₇: I=+7, F= -1
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Types of Redox Reactions
Combination Reactions: Two or more substances combine to form a single product
Example: S(s) + O₂ (g)→ SO₂ (g)
Decomposition Reaction: The breakdown of a compound into two or more components
Example: 2HgO (s) →2Hg (l) + O₂ (g)
Combustion Reaction: A substance reacts with oxygen, creating hydrocarbons, carbon dioxide, and water
Example: C₃H₈ (g) + 5O₂ (g) → 3CO₂ + 4H₂O (l)
Displacement Reaction: An ion (or atom) in a compound is replaced by an ion (or atom) of another element
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Types of Displacement Reactions
Hydrogen Displacement: Alkali metals and some alkaline earth metals (Ca, Sr, and Ba) will displace hydrogen from cold water
Example: Ca (s) + 2H₂O (l) → Ca(OH)₂ (s) + H₂ (g)
Metal Displacement: A metal compound is displaced by another metal in the elemental state
Example: TiCl₄ (g) + 2Mg (l)→Ti(s) + 2MgCl₂ (l)
Halogen Displacement: Group 7A elements can displace other group 7A elements lower on the periodic table
Example: Cl₂ (g) + 2KBr (aq) → 2KCl (aq) + Br₂ (l)
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To predict whether a metal or hydrogen displacement will occur, refer to an activity series
Any metal above hydrogen will displace it from water or acid, but metals below hydrogen will not react
Any metal listed higher than another will displace the lower metal in a metal displacement reaction
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Disproportionation Reactions
An element in one oxidation state is simultaneously oxidized and reduced
One reactant must contain an element with at least three oxidation states
Example: 2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)
In the reactants, O₂ has an oxidation number of 1-, in H₂O it is 2-, but in O₂ it is 0
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Practice Problems
Determine the type of redox reaction
Cl₂ (g) + 2NaI (aq) → 2NaCl (aq) + I₂ (s): Combination Reaction
Mg(s) + 2HCl (aq) → MgCl₂ (aq) + H₂ (g): Displacement Reaction
2KClO₃ (s) → 2KCl (s) + 3O₂ (g): Decomposition Reaction
V₂O₅ (s) + 5Ca (l) → 2V (l) + 5CaO (s): Displacement Reaction
2Al (s) + 3Br₂ (l) → 2AlBr₃ (s): Displacement Reaction
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Topic 4.8 Introduction to Acid-Base Reactions (MH 4.3)
Objective: Identify species as Bronsted-Lowry acids, bases, and/or conjugate acid-base pairs, based on proton transfer involving those species
Define Bronsted-Lowry acid and base
Water plays an important role in many acid-base reactions in aqueous solutions
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Bronsted-Lowry Acids and Bases
Bronsted Acid: Proton Donor
Bronsted Base: Proton Acceptor
Hydrochloric acid is a Bronsted Acid because it donates a proton in water
Example: HCL(aq) →H+(aq) + Cl-(aq)
HCL(aq) + H2O (l) → H2O+(aq) + Cl-(aq)
H cannot be a bare atom in an aqueous solution due to its strong attraction for the negative pole (O atom)
H3O+ is a hydronium atom, a Bronsted acid (HCl) donates a proton to a Bronsted Base (H2O)
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Conjugate Acid/Base
A pair of substances that differ in formula only by a proton (H+)
Conjugate acid is the partner of the Bronsted-Lowry Base
Conjugate base is the partner of the Bronsted-Lowry acid
All acid-base reactions involving proton transfer have two conjugate acid-base pairs
Strong acids have weak conjugate bases and vice versa
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Strong and Weak Acids
Table 4.3 Some Common Strong and Weak Acids
Strong Acids: Hydrochloric acid (HCl), Hydrobromic acid (HBr), Hydroiodic acid (HI), Nitric acid (HNO3), Sulfuric acid (H2SO4), Perchloric acid (HC1O4)
Weak Acids: Hydrofluoric acid (HF), Nitrous acid (HNO2), Phosphoric acid (H3PO4), Acetic acid (CH3COOH)
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Bronsted-Lowry Practice Problem
Identify the Bronsted Lowry acid in the following chemical equation: H3PO4(aq) + H2O (l) → H2PO4-(aq) +H3O+(aq)
Identify the Bronsted Lowry base in the following chemical equation: NH3(aq) + H2O(l) → NH4 +(aq) + OH-(aq)
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Conjugate Practice Problems
Identify the Conjugate acid and base in the following chemical equation and match it with the Bronsted Lowry acid and base:
HCL(aq) + H2O(aq) ←→ H3O+(aq) + Cl-(aq)
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H3PO4(aq) + H2O (l) → H2PO4-(aq) +H3O+(aq)
NH3(aq) + H2O(l) → NH4 +(aq) + OH-(aq)
HCL(aq) + H2O(aq) ←→ H3O+(aq) + Cl-(aq)
Orange in the reactant is the BL base and in the product is the Conjugate Acid
Blue in the reactant is the BL acid and in the product the Conjugate Base
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Acid-Base Neutralization
Any reaction between an acid and a base is called a neutralization reaction.
Aqueous neutralization reactions produce water and a salt.
A salt is an ionic compound made up of a cation other than H+ and an anion other than OH- or O2 -
Example: HCl(aq) and NaOH(aq) yield NaCl(aq) and H20 (L)
Ionic equation: H+(aq)+Cl-(aq)+Na+(aq)+OH-(aq)----->Na+(aq)+Cl-(aq)+H2O(L)
Net ionic equation: H+(aq)+OH-(aq)----->H20(L)
Na+ and Cl- are spectator ions because the charges do not change during the reaction.
This occurs because both the reactants are a strong acid and a strong base.
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Acid-Base Neutralization
This is not the case with either a weak acid or base, even if the other reactant is strong.
Example: HCN(aq)+NaOH>>>NaCN(aq)+H20(L)
Ionic equation: HCN(aq)+Na+(aq)+OH-(aq)----->Na+(aq)+CN-(aq)+H2O(L)
Net ionic equation: HCN(aq)+OH-(aq)----->CN-(aq)+H2O(L)
Only Na+ goes away because the CN- in the reactants is not by itself.
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Neutralization Leading to Gas Formation
In the equation HCl(aq)----->H+(aq)+Cl-(aq), the H+ ion is simply a bare proton.
It naturally attracts to the negative pole, or the O atom, in H2O.
This new water molecule (H3O+) is said to be hydrated.
A hydrated water molecule can have one of its hydrogens attract to a dissolved negative ion and break off from the water molecule.
In both of these cases, the water molecule is involved in the neutralization reaction.
True equation: HCl(aq)+ H2 O(L)----->H3O+(aq)+Cl-(aq)
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Practice Problems
Find a) the ionic equation and b) the net ionic equation of each of the following:
2HBr(aq)+Ba(OH)2(aq)----->BaBr2(aq)+2H2O(L)
H2SO4(aq)+2KOH(aq)----->K2SO4(aq)+2H2O(L)
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Answers
2H+(aq)+2Br-(aq)+Ba2+(aq)+2OH-(aq)----->Ba2+(aq)+2Br-(aq)+2H2O(L)
2H+(aq)+2OH-(aq)----->2H2O(L)
H+(aq)+ HSO4-(aq)+ 2K+(aq)+2OH-(aq)----->2K+(aq)+SO42 -(aq)+ 2H2O(L)
H+(aq)+HSO4-(aq)+2OH-(aq)----->SO42 -(aq)+2H2O(L)
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Bases be like: I am once again asking for your Protons
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reactants products
Call an ambulance!
products reactants
But not for me!
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11:20 LTE Evil Magnesium be like
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11:17 LTE Things I learned in Organic Chemistry
Interesting Reactions
Deadly compounds
Nomenclature
How to draw hexagons
SOUTH memes
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11:15 LTE P-CHEM FANS hahaha
THSINOTCYCLOH EXA-1.0.5-TRIENE. noo000000000
ITSIBENZENE BIOLOGY ENJOYERS
THIS BIOCHEMIIS DUDE.
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11:26 LTE when you fail your chemistry exam but can still understand chem memes
I'm still worthy!
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Topic 4.9 Oxidation-Reduction reactions (MH 4.4)
Objective: represent a balanced redox reaction equation using half reactions.
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Oxidation-Reduction Reactions
Electron transfer (charges change)
Oxidation is loss of electrons
Reduction is a gain of electrons
Example: Mg + Mg2+ O2
Oxidation half-reaction: 2Mg 2Mg2+ + 4e-
Reduction half-reaction: O2 + 4e- 202-
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Oxidation Number
The charge the atom would have in a molecule if electrons were completely transferred.
Rules:
Free elements (uncombined state) have an oxidation number of zero.
In monatomic ions, the oxidation number is equal to the charge on the ion.
The oxidation number of oxygen is usually -2. In peroxides, it is -1.
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Oxidation Numbering Rules continued.
The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is -1.
Group IA metals are +1, IIA metals are +2 and fluorine is always -1.
The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. (neutral molecules = 0)
Oxidation numbers do not have to be integers. The oxidation number of oxygen in the superoxide ion, O2- is -½
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Example 5
Assign oxidation numbers to all the elements in the following compounds and ion:
Li2O
HNO3
Cr2O72 -
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Types of Oxidation-Reduction Reactions
Displacement Reaction: A+BC→AC +B
Combination Reaction: A +B →C
Decomposition Reaction: C →A+B
Combustion Reactions: A + O2 →B
Disproportionation Reaction: the same element is simultaneously oxidized and reduced.
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Example 6
Classify the following redox reactions and indicate changes in the oxidation numbers of the elements:
(a) 2N20(g) 2N2(g)+O2(g)
(b) 6Li(s) + N2(g) 2Li3N(s)
(c) Ni(s) + Pb(N03)2(aq) Pb(s) + Ni(NO3)2(aq)
(d) 2NO(g) + H20(1) HNO(aq) + HNO3(aq)
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Electrochemical Processes
Oxidation-reduction reactions in which:
The energy released by a spontaneous reaction is converted into electricity or
Electrical energy is used to cause a nonspontaneous reaction to occur.
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Half-Reactions
Oxidation reactions can be considered as two separate processes occurring simultaneously.
The whole reaction can be thought of as the sum of two halves-an oxidation half and a reduction half.
Oxidation half reactions are always written with the electrons on the right side of the equation (as products after the loss)
Reduction half reactions are always written with the electrons on the left side of the equation (as reactants taken by the original element)
An ionic equation needs to be written before attempting to write the half reactions equations.
Example: AgNO3(aq) + Cu (s) —--> Cu(NO3)2 (aq)+ 2 Ag (s)
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Balancing Redox Equations
Write the unbalanced equation for the reaction in ionic form.
Separate the equation into two half-reactions.
Balance the atoms other than O and H in each half-reaction.
For reactions in acid, add H2O to balance Oxygen atoms and H+ to balance the hydrogen atoms.
Add electrons to one side of each half reaction to balance the charges on the half-reactions.
If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by correct coefficients.
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Balancing Redox Equations cont. 7. Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. 8. Verify that the number of atoms and the charges are balanced. 9. For reactions that are in basic solutions, add OH- to both sides of the equation for every H+ that appears in